CHEMISTRY S2

Comprehensive

Chemistry

for Rwanda Schools

Student’s Book

Secondary 2

Authors

Mukesha Sandra

Mukazi Ndekezi

Thomas Grissell

Edward Kemp

Laxmi Publications (P) Ltd

(An ISO 9001:2008 Company)

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Comprehensive CHEMISTRY FOR RWANDA SCHOOLS STUDENT’S BOOK—SECONDARY 2

Published by

Laxmi Publications Pvt. Ltd

113, Golden House, Daryaganj,

New Delhi-110002, India

Telephone: 91-11-4353 2500, 4353 2501

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Rwanda Address:

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© All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form by any means, electronic or mechanical, photocopying, recording or otherwise, without the prior permission of publishers.

First published, 2017

ISBN: 978-93-5274-031-4

Printed and bound in India

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Contents

Preface (vii)

Unit 1. Chemical Bonding........................................................................................................ 1–21

1.1 Stability of Atoms......................................................................................................................... 1

1.2 Formation of Ions from Atoms.............................................................................................. 3

1.3 Ionic Bonding.................................................................................................................................. 5

1.4 Formation of Ionic Bond........................................................................................................... 6

1.5 Properties of Ionic Compounds.............................................................................................. 7

1.6 Covalent Bonding.......................................................................................................................... 8

1.7 Formation of Covalent Bond................................................................................................... 9

1.8 Properties of Covalent Compounds..................................................................................... 9

1.9 Giant Covalent Structures....................................................................................................... 11

1.10 Metallic Bonding......................................................................................................................... 14

1.11 Formation of Metallic Bond.................................................................................................. 14

1.12 Properties of Metallic Bond................................................................................................... 15

1.13 Summary......................................................................................................................................... 17

1.14 Glossary........................................................................................................................................... 18

1.15 Unit Assessment.......................................................................................................................... 18

Unit 2. Trends in Properties of Elements in the

Periodic Table 22–53

2.1 Classification of Elements...................................................................................................... 22

2.2 Physical Properties of Metals................................................................................................ 28

2.3 Physical Properties of Non-metals...................................................................................... 34

2.4 Trends in Reactivity for Metals and Non-metals........................................................ 36

2.5 Chemical Properties of Metals............................................................................................. 37

2.6 Chemical Properties of Non-metals................................................................................... 46

2.7 Comparison Among the Physical and Chemical Properties of Metals

and Non-metals 47

2.8 Uses of Metals and Non-metals........................................................................................... 48

2.9 Summary......................................................................................................................................... 49

2.10 Glossary........................................................................................................................................... 50

2.11 Unit Assessment.......................................................................................................................... 50

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Unit 3. Water Pollution............................................................................................................ 54–71

3.1 Water Pollution............................................................................................................................ 54

3.2 Main Water Pollutants............................................................................................................. 56

3.3 Dangers of Polluted Water..................................................................................................... 61

3.4 Prevention of Water Pollution.............................................................................................. 65

3.5 Our Clean Future......................................................................................................................... 67

3.6 Summary......................................................................................................................................... 68

3.7 Glossary........................................................................................................................................... 69

3.8 Unit Assessment.......................................................................................................................... 70

Unit 4. Effective Ways of Waste Management................................................ 72–92

4.1 Steps to Effective Waste Management............................................................................ 72

4.2 Importance and Benefits of Waste Recycling............................................................... 82

4.3 Effects of Waste and Poor Disposal.................................................................................. 87

4.4 Summary......................................................................................................................................... 90

4.5 Glossary........................................................................................................................................... 91

4.6 Unit Assessment.......................................................................................................................... 91

Unit 5. Categories of Chemical Reactions........................................................ 93–126

5.1 Types of Reactions..................................................................................................................... 94

5.2 Classification of Chemical Reactions as Endothermic and

Exothermic Reactions............................................................................................................ 114

5.3 Ionic Equations.......................................................................................................................... 118

5.4 Rules for Writing Ionic Equations................................................................................... 118

5.5 Summary....................................................................................................................................... 121

5.6 Glossary......................................................................................................................................... 122

5.7 Unit Assessment....................................................................................................................... 123

Unit 6. Preparation of Salts and Identification of Ions.................... 127–155

6.1 Saturated and Unsaturated Solutions.............................................................................. 129

6.2 Supersaturated Solutions....................................................................................................... 130

6.3 Factors Influencing Solubility of Different Salts....................................................... 131

6.4 Solubility Curve......................................................................................................................... 133

6.5 Calculation of Solubility........................................................................................................ 137

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6.6 Different Ways of Preparing Normal Salts.................................................................. 140

6.7 Uses and Sources of Salt....................................................................................................... 144

6.8 Identification of Ions.............................................................................................................. 146

6.9 Summary....................................................................................................................................... 151

6.10 Glossary......................................................................................................................................... 152

6.11 Unit Assessment....................................................................................................................... 153

Unit 7. The Mole Concept and Gas Laws..................................................... 156–210

7.1 Avogadro Number and the Mole Concept.................................................................. 156

7.2 Calculation of the Number of Moles.............................................................................. 160

7.3 Definition of Relative Atomic Mass............................................................................... 161

7.4 Definition and Calculation of Relative Molecular Mass....................................... 162

7.5 Definition and Calculation of Relative Formula Mass.......................................... 163

7.6 Calculation of Molar Mass.................................................................................................. 164

7.7 Relationship between Number of Moles, Mass and Molar Mass.................... 165

7.8 Calculation of Mass Percent Composition of an Element in a Compound 174

7.9 Empirical Formula and Molecular Formula................................................................ 178

7.10 Stoichiometric Calculations................................................................................................. 186

7.11 Limiting Reactants................................................................................................................... 189

7.12 The Gaseous State.................................................................................................................... 191

7.13 The Gas Laws............................................................................................................................. 192

7.14 Calculation of Molar Gas Volume under Standard Conditions......................... 206

7.15 Summary....................................................................................................................................... 207

7.16 Glossary......................................................................................................................................... 208

7.17 Unit Assessment....................................................................................................................... 208

Unit 8. Preparation and Classification of Oxides.................................. 211–228

8.1 Preparation of Oxides............................................................................................................. 211

8.2 Reaction of Oxides with Water......................................................................................... 218

8.3 Reaction of Oxides with Acids and Bases.................................................................... 220

8.4 Classification of Oxides......................................................................................................... 221

8.5 Uses and Production of Slaked Lime (ISHWAGARA)........................................... 225

8.6 Summary....................................................................................................................................... 226

8.7 Glossary......................................................................................................................................... 226

8.8 Unit Assessment....................................................................................................................... 226

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Unit 9. Electrolytes and Non-electrolytes....................................................... 229–238

9.1 Electrolyte and Non-electrolyte......................................................................................... 229

9.2 Definition of Electrolysis...................................................................................................... 230

9.3 Strong Electrolyte..................................................................................................................... 230

9.4 Weak Electrolyte...................................................................................................................... 231

9.5 Conductivity of Electricity by Electrolytes................................................................. 233

9.6 Application of Electrolytes.................................................................................................. 234

9.7 Summary....................................................................................................................................... 235

9.8 Glossary......................................................................................................................................... 235

9.9 Unit Assessment....................................................................................................................... 236

Unit 10. Properties of Organic Compounds and

Uses of Alkanes 239–254

10.1 Organic Chemistry................................................................................................................... 240

10.2 Difference between Organic and Inorganic Chemistry.......................................... 240

10.3 Occurrence of Organic Compounds................................................................................ 242

10.4 Homologous Series.................................................................................................................. 242

10.5 General Formula of Alkanes............................................................................................... 243

10.6 Nomenclature of Alkanes..................................................................................................... 243

10.7 Structural Isomerism............................................................................................................... 244

10.8 Physical Properties of Alkanes........................................................................................... 245

10.9 Chemical Properties of Alkanes........................................................................................ 247

10.10 Laboratory Preparation of Methane................................................................................ 248

10.11 Uses of Alkanes......................................................................................................................... 249

10.12 Summary....................................................................................................................................... 250

10.13 Glossary......................................................................................................................................... 251

10.14 Unit Assessment....................................................................................................................... 252

Bibiliography....................................................................................................................................... 255

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Preface

Our aim was to write a book that should be sensitive to the needs of learners and help them build relevant skills and competences that prepare them to be well integrated in the society and exploit employment opportunities. Learners in the 21^st century have to overcome the challenges brought about by the rapid developments in science, technology and society. This book provides them with learning experiences that enable them to attain knowledge and develop a global perspective and life-long learning skills, so that they can contribute to today’s knowledge-based economy and society.

This book, Comprehensive Chemistry, has been developed on the basis of the new Ordinary Level Chemistry syllabus in accordance with the recommendations of Rwanda Education Board (REB). It provides skills to the learners that guide the construction of theories and laws that help to explain natural phenomenon and manage people and the environment. The units covered in this book help learners to have a better understanding of the impact and influence chemistry has in a modern scientific world. The units include chemical bonding; trends in properties of elements in the periodic table; water pollution; effective ways of waste management; categories of chemical reactions; preparation of salts and identification of ions; the mole concept and gas laws; preparation and classification of oxides; electrolytes and non-electrolytes and properties of organic compounds and uses of alkanes.

Illustrations and activities given in the text help learners to obtain practical learning. Through experimentation, observation and presentation of information during the learning process, the learners develop not only deductive and inductive skills but also communication, critical thinking and problem solving skills as they try to make inferences and conclusions. Assessment at the end of each unit will help them evaluate their skills. Glossary given at the end of each unit will help the learners know the meaning of difficult words.

This book develops a sense of national identity and teaches learners to be committed to contribute to the economic development of the nation and the world. It enhances curiosity and interest in the natural and technological world as well as understanding the impact of science and technology on society. An honest attempt has been made to develop care and concern for the environment.

Any suggestions for the improvement of this book will be gratefully appreciated and acknowledged.

—Authors

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Unit 1

Chemical Bonding

LEARNING OBJECTIVES KNOWLEDGE GAIN

After reading this unit, you will be able to:

•• explain the nature of ionic, covalent and metallic bonding.

•• state the typical physical properties of ionic compounds, and of covalent compounds.

•• explain the physical properties of metals in terms of their structure.

In 1985, a new allotrope of carbon Buckminsterfullerene was discovered. It has a cage-like ring structure which resembles a football. It is made of twenty hexagons and twelve pentagons.

1.1 STABILITY OF ATOMS

ACTIVITY 1.1: Showing Stability of Atoms

•• Take a glass full of water. Try adding water into it. Are you able to add? •• Now take another glass of water but a quarter (one-fourth) filled. •• Try adding water into it. Now, are you able to add or not?

Perform the two activities in classroom and then discuss your answers among your classmates.

In the above activity, you will observe that when the glass was already filled, there was no space to add more water into it. Thus, the water in the glass remained stable. A noble gas has a fully filled outermost shell just like the glass full of water. It has eight electrons in the outermost shells except helium (2 electrons).

When atoms or the elements combine to form molecules, a force of attraction is developed between the atoms (or ions) which holds them together. The force which links the atoms (or ions) in a compound is called a chemical bond (or just “bond”). A bond is formed so that each atom acquires a stable electronic configuration similar to that of a noble gas.

1

Table 1.1: Electronic Configurations of Noble Gases (or Inert Gases)

Noble gas Symbol Atomic Electronic configuration Number of electrons in
(Inert gas) number K L M N O P outermost shell

Helium He 2 2 2

Neon Ne 10 2, 8 8

Argon Ar 18 2, 8, 8 8

Krypton Kr 36 2, 8, 18, 8 8

Xenon Xe 54 2, 8, 18, 18, 8 8

Radon Rn 86 2, 8, 18, 32, 18, 8 8

The atoms combine with one another to achieve the inert gas electron arrangement and become more stable. So, when atoms combine to form compounds, they do so in such a way that each atom gets 8 electrons in its outermost shell or 2 electrons in the outermost n shell.

An atom can achieve the inert gas electron arrangement in three ways:

•• By losing one or more electrons (to another atom).

Atoms with 1, 2 or 3 electrons in the outermost shell lose electrons to achieve stability.

•• By gaining one or more electrons (from another atom).

Atoms with five, six or seven electrons in the outermost shell gain three, two or one electron respectively to achieve stability.

•• By sharing one or more electrons (with another atom).

Atoms with four to seven electrons in outermost shell may achieve stability by sharing them with each other.

Table 1.2: Electronic Configuration of Some Metals and Non-metals

Type of element Element Atomic Number of electrons in shells
number K L M N

Metals Sodium (Na) 11 2 8 1
Magnesium (Mg) 12 2 8 2
Aluminium (Al) 13 2 8 3
Potassium (K) 19 2 8 8 1
Calcium (Ca) 20 2 8 8 2
Non-metals Nitrogen (N) 7 2 5
Oxygen (O) 8 2 6
Fluorine (F) 9 2 7
Phosphorus (P) 15 2 8 5
Sulphur (S) 16 2 8 6
Chlorine (Cl) 17 2 8 7

2 S2 Chemistry

EXERCISE 1.1

What do you mean by a chemical bond?

When a bond is formed, each atom acquires a stable configuration similar to _______

Generally, metals lose electrons to achieve inert gas electron arrangement.
(True or False)

Which of the following is not a noble gas?

(a) Helium (b) Neon

(c) Hydrogen (d) Argon

Among, phosphorus, sulphur, and calcium; which element achieves stability by losing electron.

1.2 FORMATION OF IONS FROM ATOMS

ACTIVITY 1.2: Illustrating Formation of Ion

Divide the class into two groups. Half of the students hold positive plank cards and another half hold negative plank cards. Positive plank cards are protons and negative plank cards are electrons. Now perform the following and analyse:

Students with 5 positive and 5 negative plank cards are grouped together. Their total charge being neutral in the group.

•• Now, one electron is removed from the group. 4 students are left holding negative plank cards.

Can you tell the net charge now of this group?

•• Add one electron to the neutral group. 6 students are now holding negative plank cards.

Can you now tell what is the charge of this group?
•• Similarly, perform the above activity with 7 students and analyse the charge.

An atom contains electrons, protons and neutrons.

Protons carry positive charges, electrons carry negative charges and neutrons carry no charges. Every atom contains an equal number of “positively charged protons” and “negatively charged electrons”. Thus, an atom is electrically neutral.

An ion is formed when an atom loses or gains one or more electrons. The atom may be of a metal or a non-metal.

A metal readily loses its outermost electron or electrons to form a positive ion or cation. The number of positive charges carried on a cation is equal to the number of electron(s) lost by the metal atom. Examples are given in Table 1.3.

Table 1.3: Some Metals and their Ions

Parent atom Electrons Name and
(electronic lost symbol of ion
configuration) (electronic
configuration)

Sodium 1 Sodium ion,
(2, 8, 1) Na+ (2, 8)

Calcium 2 Calcium ion,
(2, 8, 8, 2) Ca2+ (2, 8, 8)

Aluminium 3 Aluminium
(2, 8, 3) ion, Al3+ (2, 8)

Chemical Bonding 3

Metal ions carry positive charges because the number of positively charged protons in the nucleus becomes greater than the number of negatively charged electrons surrounding it. For example, in a sodium atom there are 11 protons in the nucleus and 11 electrons surrounding it. Loss of one electron to form a sodium ion means that there are 11 protons but only 10 electrons. There is a net charge of 1+. This charge is written as a superscript at the right of the symbol of the element (Figure 1.1).

Loses this 1 electron
M
L
The last shell 11p K –1 electron,

disappears 12n

+

L
11p K
12n

Sodium atom (Na) Sodium ion (Na+)

Figure 1.1: Formation of a sodium ion.

Hydrogen atoms can also lose an electron to form an ion with one positive charge.

Some non-metals readily gain one or more electrons into their outermost shell to form a negative ion or anion. The number of negative charges an anion carries is equal to the number of electron(s) gained by the non-metal atom. Examples are given in Table 1.4.

Table 1.4: Some Non-metals and their Ions

Parent atom Electrons gained Name and symbol of ion
(electronic configuration) (electronic configuration)

Chlorine (2, 8, 7) 1 Chloride ion, Cl– (2, 8, 8)
Oxygen (2, 6) 2 Oxide ion, O2– (2, 8)
Nitrogen (2, 5) 3 Nitride ion, N3– (2, 8)

Non-metal ions carry negative charges because the number of negatively charged electrons surrounding the nucleus becomes greater than the number of positively charged protons in it. For example, in a chlorine atom there are 17 protons in the nucleus and 17 electrons surrounding it. Gain of one electron to form a chloride ion means that there are 18 electrons and only 17 protons. There is a net charge of –1. This charge is written as a superscript at the right of the symbol of the element (Figure 1.2).

4 S2 Chemistry

Gains this 1 electron

M
L
+1 electron,
11p K
18n

–

M

L
K
11p
18n

Chlorine atom (Cl) Chloride ion (Cl–)

Figure 1.2: Formation of a chloride ion.

Notice that when a non-metal forms an anion, the name changes slightly; chlorine forms a chloride ion, oxygen forms an oxide ion. Several common radicals exist as negative ions including nitrate (NO–3), carbonate (CO23– ) and phosphate (PO43–).

EXERCISE 1.2

Define ion.

Which of the following is an anion?
(a) Cl– (b) Na+
(c) Mg2+ (d) Al3+
Why do metal ions carry positive charges?

The number of negative charges an anion carries is equal to the number of electrons gained by the non-metal atom. (True or False)
Give two examples of each:

(i) anion (ii ) cation

1.3 IONIC BONDING

The compounds which are made up of ions are known as ionic compounds. In an ionic compound, the positively charged ions (cations) and negatively charged ions (anions) are held together by the strong electrostatic forces of attraction. The forces which hold the ions together in an ionic compound are known as ionic bonds or electrovalent bonds. Since an ionic bond consists of an equal number of positive ions and negative ions, the overall charge on an ionic compound is zero. For example, sodium chloride (NaCl) is an ionic compound which is made up of equal number of positively charged sodium ions (Na+) and negatively charged chloride ions (Cl–). Some of the common ionic compounds, their formulae and the ions present in them are given in Table 1.5.

Table 1.5: Formulae and Nomenclature of Some Ionic Compounds

Nomenclature Formula Ions present
Aluminium oxide Al2O3 Al3+ and O2–
Ammonium chloride NH Cl NH + and Cl–
4 4
Calcium hydroxide Ca(OH) 2 Ca2+ and OH–

Chemical Bonding 5

Calcium nitrate
Ca(NO 3 ) 2 Ca2+ and NO –
3
Calcium oxide CaO Ca2+ and O2–
Copper sulphate CuSO Cu2+ and SO 2–
4 4
Magnesium chloride MgCl 2 Mg2+ and Cl–

Potassium chloride KCl K+ and Cl–
Potassium hydroxide KOH K+ and OH–
Sodium carbonate Na CO 3 Na+ and CO2–
2 3
Sodium hydroxide NaOH Na+ and OH–

Ionic compounds are made up of a metal and a non-metal (except ammonium chloride which is an ionic compound made up of only non-metals). So, whenever a bond involves a metal and a non-metal, we call it ionic bond.

resulting in formation of sodium chloride.
Na Na+ + e–
2,8,1 2, 8

Sodium ion
(Cation)

EXERCISE 1.3

Give two examples of ionic compounds. Write their chemical formulae.
The overall charge on an ionic compound is zero. (True or False)

Name the ions present in calcium nitrate.

Ionic compounds are made up of a

______ and a ______ .

Give an example of an ionic compound made up of only non-metals.

1.4 FORMATION OF IONIC BOND

An ionic bond changes the electronic configurations of the atoms. Metal atoms lose their outermost electron(s), forming cations. Non-metal atoms gain electron(s) to fill their outermost shell, forming anions. The electrostatic force of attraction between the oppositely charged ions holds the ions together. For example,

(a) When a hot sodium atom is placed in chlorine gas, a reaction takes place

S2 Chemistry

Cl + e– Cl–
2,8,7 2,8,8
Chloride ion
(Anion)
. ×× (Na+) ××–

Na + ×Cl×× ×.Cl××
2,8,1 ×× ××

Formation of Sodium Chloride

(b) When a magnesium atom comes in contact with chlorine gas, it forms magnesium chloride.

Mg Mg2+ + 2e–
2,8,2 2, 8
Magnesium ion
(Cation)
Cl + e– Cl–
2,8,7 2,8,8
Chloride ion
(Anion)
××
×Cl×× ×× –
Mg: + [Mg2+ ] .
×× ×Cl××
×× × ×× 2
2,8,1
×Cl×
××

Formation of Magnesium Chloride

EXERCISE 1.4

With the help of dot and cross, show the formation of CaCl2.

Which of the following is correct?
– –
(a) [Na + ] . ×× × – ] Cl××
× Cl××× (b) [Na . ××
+ –
(c) [Na+] ×× (d ) [Na]+ ××Cl××
××Cl ××
×× ××
3. The electrostatic force of attraction
holds the ions together.
(True or False)

1.5 PROPERTIES OF IONIC COMPOUNDS

ACTIVITY 1.3: Illustrating

Physical Properties of Ionic

Compounds

•• Take a sample of sodium chloride or any other salt from the science laboratory.
•• What is the physical state of this salt?

•• Take a small amount of a sample on a metal spatula and heat directly on the flame as shown in figure (a).

Spatula
containing sample

Bunsen burner

Testing melting point of sodium chloride

•• What did you observe? Did the sample impart any colour to the flame? Does this compound melt?

•• Try to dissolve the sample in water, petrol and kerosene. Is it soluble?

•• Make a circuit as shown in figure (b) and insert the electrodes into a solution of salt. What did you observe?
Battery Bulb

Switch

Beaker

Graphite rod

Salt solution

under test

Testing electrical conductivity of salt solution

•• What is your inference about the nature of this compound?

You may have observed the following general properties for ionic compounds:

•• Ionic compounds are usually crystalline solids.

•• Ionic compounds have high melting and high boiling points.

The temperature at which a solid melts into liquids is called the melting point of the solid. The temperature does not change during melting.

Boiling point is the temperature at which a liquid changes into a gas. The temperature of a liquid remains the same once boiling has started.

Chemical Bonding 7

Table 1.6: Melting and Boiling Points of

Some Ionic Compounds

Ionic Melting Boiling point
compounds point (K) (K)

NaCl 1074 1686
LiCl 887 1600
CaCl2 1045 1900
CaO 2850 3120
MgCl2 981 1685

•• Ionic compounds are usually soluble in water but insoluble in organic solvents like petrol and kerosene.

•• Ionic compounds conduct electricity when dissolved in water or when melted. When we dissolve the ionic solid in water

1.6 COVALENT BONDING

or melt it, the crystal structure is broken down to form ions. These ions help in conducting electricity.

EXERCISE 1.5

Why do ionic compounds conduct electricity when dissolved in water?

Ionic compounds are insoluble in

(a) kerosene

(b) petrol

(c) both (a) and (b)

(d) neither (a) nor (b)

Ionic compounds have low melting and boiling points. (True or False)

Ionic compounds are usually ______

solids.

The chemical bond formed by sharing of electrons between two atoms is known as a covalent bond. The compounds containing covalent bonds are known as covalent compounds. A covalent bond is formed when both the reacting atoms need electrons to achieve the inert gas electron arrangement. Now, the non-metals have usually 5, 6 or 7 electrons in the outermost shells of their atoms. So, all the non-metal atoms need electrons to achieve the inert gas structure. They get these electrons by mutual sharing. Thus, whenever a non-metal combines with another non-metal, covalent bond is formed.

Table 1.7: Formulae and Nomenclature of Some Covalent Compounds

Nomenclature Formula Elements present

Methane CH4 C and H
Ethane C2H6 C and H
Ethene C2H4 C and H
Ethyne C2H2 C and H
Ammonia NH3 N and H
Alcohol (Ethyl alcohol) C2H5OH C, H and O
Hydrogen sulphide gas H2S H and S
Carbon dioxide CO2 C and O
Carbon disulphide CS2 C and S

8 S2 Chemistry

Carbon tetrachloride
CCl4 C and Cl
Glucose C6H12O6 C, H and O
Cane sugar C12 H22O11 C, H and O
Urea CO(NH2)2 C, O, N and H

EXERCISE 1.6

What do you mean by a covalent bond?

Give two examples of covalent compounds. Also write their chemical formulae.

When a ______ combines with another ______, covalent bond is formed.

Choose the covalent compound(s).
(a) CH4 (b) H2S (c) CS2 (d) All of these

1.7 FORMATION OF COVALENT BOND

Covalent bonding between atoms of different elements.

(i) Carbon atom shares four electrons to form methane.

. ×H H
.
×H ×
.C. + H×. C .× H
. ×H ×.
×H
Carbon Hydrogen H

atom atoms Methane

Covalent bonding or sharing of electrons only takes place in outermost shells of atoms to attain inert gas electron arrangement.

(ii) As in water molecule, 2 hydrogen atoms share electrons with oxygen atom.

: ×H H ×. :
.O. + O×. H
: ×H :
Oxygen Hydrogen Water molecule
atom atoms
H–O–H or H2O

EXERCISE 1.7

Using cross and dot diagram, show the formation of carbon dioxide.
Which of the following is correct?
(a) (b)

(c) (d)

Carbon tetrachloride and Urea are not covalent compounds. (True or False)
In the formation of covalent bonding,

______ of electrons takes place in the

______ shells of atoms.

1.8 PROPERTIES OF COVALENT COMPOUNDS

ACTIVITY 1.4: Illustrating Physical Properties of Covalent Compounds

Let us test some covalent compounds in different ways:

•• Take sample of cooking oil. Try to dissolve it in water and ethanol. Does it dissolve?

Chemical Bonding 9

•• Have you ever observed a burning candle wax? If not, take a candle wax and observe it burning. How much time does it take to melt down?

•• Take a pan and add water to it. Let it boil. Do you know the boiling point of water?
•• Now add two electrodes to the water pan making a circuit. What did you observe? What would have happened if you would have added NaCl salt in the pan?

•• What can you now say about these covalent compounds?

You have observed the following properties of covalent compounds:

•• Covalent compounds are usually liquids, gases or solids. For example, alcohol, benzene, water and cooking oil are liquids. Methane, ethane and chlorine are gases. Glucose, urea, and wax are solid covalent compounds.

•• Covalent compounds have usually low melting points and low boiling points.

•• Covalent compounds are usually insoluble in water, but they are soluble in organic solvents. Some of the covalent compounds like glucose, sugar and urea are soluble in water.

•• Covalent compounds do not conduct electricity because they do not contain ions.

ACTIVITY 1.5: Detecting an

Ionic Bond or Covalent Bond

•• Take the sample such as common salt (NaCl) provided.
•• Try to dissolve it in water.

•• If it dissolves, chances are it is likely to be an ionic compound. But, you already know some covalent compounds like glucose, urea and sugar are soluble in water.

•• Now, perform electrical conductivity test.

•• If the NaCl sample dissolves in water, arrange a circuit with two electrodes and a bulb.
•• Figure out whether the bulb glows or not. According to your observation conclude the bond present in the sample.
•• Make a report on the properties of ionic and covalent compounds.

Bulb

Battery

Electrode

Common
Glucose salt

Table 1.8: Differences between Ionic Compounds and Covalent Compounds

Ionic compounds Covalent compounds

Ionic compounds are usually crystalline solids. Covalent compounds are usually solids, liquids

or gases.

Ionic compounds have high melting points and Covalent compounds have usually low melting

boiling points. That is, ionic compounds are and boiling points.
non-volatile.
Ionic compounds conduct electricity when Most covalent compounds do not conduct
dissolved in water or melted. electricity.

10 S2 Chemistry

Ionic compounds are usually soluble in water.
Covalent compounds are usually insoluble in
water (except, glucose, sugar, urea, etc.).

Ionic compounds are insoluble in organic Covalent compounds are soluble in organic
solvents (like alcohol, ether, acetone, etc.) solvents.

ACTIVITY 1.6: Identifying Ionic and Covalent Compounds

Choose the ionic as well as covalent compounds from the bubbles and make a table in your exercise notebook.

Nitric acid
Methane Urea Ionic
compound

H2O Acetic acid
Ethyl CuSO4
alcohol
Glucose Formaldehyde Covalent
compound
H2SO4 CaO NaOH

EXERCISE 1.8

Some covalent compounds are solid. (True or False)

Most covalent compounds are ______

in water but ______ in organic solvents.

Name two covalent compounds which are soluble in water.

Why most covalent compounds do not conduct electricity?

Melting and boiling points of covalent compounds are

(a) high

(b) low

(c) between 500°C and 1000°C

(d) cannot be determined.

1.9 GIANT COVALENT

STRUCTURES

Diamond, graphite and silicon dioxide have giant covalent structures.

1.9.1 Diamond and its Properties

Diamond is a colourless transparent substance having extraordinary brilliance. Diamond is quite heavy. Diamond is extremely hard. It is the hardest natural substance known. Diamond does not conduct electricity. Diamond burns on strong heating to form carbon dioxide. It has a very high melting point. If we burn diamond in oxygen, then only carbon dioxide gas is formed and nothing is left behind. This shows that diamond is made up of carbon only. Since diamond is made up of carbon atoms only, its symbol is taken to be C.

Strong bonds

exist between

all the carbon

atoms

Figure 1.3: Structure of diamond (The black balls represent carbon atoms).

Chemical Bonding 11

1.9.2 Graphite and its Properties

Graphite is a greyish-black opaque substance. Graphite is lighter than diamond. Graphite is soft and slippery to touch. Graphite conducts electricity. Graphite burns on strong heating to form carbon dioxide. Like diamond, graphite also has very high melting point. If we burn graphite in oxygen, then only carbon dioxide gas is formed and nothing is left behind. This shows that graphite is made up of carbon

only. Since graphite is made up of carbon atoms only, its symbol is taken to be C.

result from the very strong covalent bonds that hold the silicon and oxygen atoms in the giant covalent structure. Silicon dioxide is found as quartz in granite, and is the major compound in sandstone. The sand on a beach is made mostly of silicon dioxide.

= Sillicon atom

= Oxygen atom

Weak forces

hold the layers

of carbon atoms

together

Strong bonds

exist between

carbon atoms

in a layer

Figure 1.4: Structure of graphite (The black balls represent carbon atoms).

1.9.3 Silicon Dioxide and its Properties

Silicon dioxide (also known as Silica) has a giant covalent structure. Each silicon atom is covalently bonded to four oxygen atoms. Each oxygen atom is covalently bonded to two silicon atoms. This means that, overall, the ratio is two oxygen atoms to each silicon atom, giving the formula SiO2. Silicon dioxide is very hard. It has a very high melting point (1,610°C) and boiling point (2,230°C). It is insoluble in water, and does not conduct electricity. These properties

S2 Chemistry

Figure 1.5: Structure of silicon dioxide

1.9.4 Uses of Diamond, Graphite and Silicon Dioxide

Uses of Diamond

•• Since diamond is extremely hard, it is used for cutting and grinding other hard materials. It is also used for drilling holes in the earth’s rocky layers. Diamond ‘dies’ are used for drawing thin wires like the tungsten filament of an electric bulb.

•• Diamonds are used for making jewellery. The use of diamonds in making jewellery is because of their extraordinary brilliance. Diamond is also used in the tip of glass cutter. A sharp diamond-edged knife called keratome is used by eye surgeons to remove cataract from the eyes.

Glass Cutter Keratome Jewellery

Figure 1.6: Some of the uses of diamond.

Uses of Graphite

Due to its softness, powdered graphite is used as a lubricant for fast moving parts of machinery. Graphite can be used as a dry lubricant in the form of graphite powder or mixed with petroleum jelly to form graphite grease. Graphite powder can also be mixed with lubricant oils.

Anode

(Zinc Inner Case)

Cathode

(Graphite Rod)

Paste of MnO2,
NH4Cl, and Carbon

Figure 1.7: Some of the uses of graphite.

•• Graphite is a good conductor of electricity due to which graphite is used for making carbon electrodes or graphite electrodes in dry cells and electric arcs. The black coloured ‘anode’ of a dry cell is made of graphite. The carbon brushes of electric motors are also made of graphite.

•• Graphite is used for making the cores of our pencils called ‘pencil leads’ and black points. Graphite is black in colour and quite soft. So, it marks black lines on paper. Due to this property, graphite is used for making pencil leads. For making pencil leads, graphite is usually mixed with clay.

Chemical Bonding 13

Uses of Silicon Dioxide

•• An estimated 95% of silicon dioxide produced is consumed in the construction industry, e.g. for the production of Portland cement

•• Silica is used primarily in the production of glass for windows, drinking glasses, beverage bottles, and many other uses.

•• The majority of optical fibres for telecommunication are also made from silica.

Glass made from silicon dioxide (Silica)

Bundle of optical fibres composed of high

purity silica

Figure 1.8: Some of the uses of silicon dioxide

S2 Chemistry

EXERCISE 1.9

Diamond and Graphite are two common allotropes of Carbon. (True or False)
Which of the following is correct?

(a) Diamond is the hardest substance known.

(b) Graphite has very low melting point.

(c) Graphite does not conduct electricity.

(d ) Diamond burns on strong heating to form helium gas.

Why are diamonds used for making jewellery?

Graphite is used for making

(a) pencil lead (b) electrodes

(c) both (a) and (b)

(d ) none of these.

Diamond and Graphite have very

______ melting point.

1.10 METALLIC BONDING

The force which binds various metal atoms together is called metallic bond. The metallic bond is neither a covalent bond nor an ionic bond because these bonds are not able to explain properties of metals.

For example, metals are very good conductors of electricity but in solid state. Both ionic and covalent compounds cannot do so with the exception of graphite.

1.11 FORMATION OF METALLIC BOND

Loreutz proposed the theory of electron gas model or electron sea model for metallic bonding.

In this model, the metal is pictured as an array of metal cations in a “sea” of electrons. The atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations. Delocalised electrons are not held by any specific atom and can move easily throughout the solid. A metallic bond is the attraction between these electrons and the metallic cation.

Positive ions

from the metal

Electron cloud that does not

belong to any one metal ion

Figure 1.9: Formation of metallic bond.

EXERCISE 1.10

Name the scientist who proposed the theory of electron sea model.

Metallic bond is neither a covalent bond nor an electrovalent bond. (True or False)

The force which binds various metal atoms together is called ______ .

Make a 3D structure of electron sea model.

Write a short note on formation of metallic bonding.

1.12 PROPERTIES OF METALLIC BOND

ACTIVITY 1.7: Illustrating the Properties of Metals

•• Take samples of iron, copper, aluminium, sodium, carbon and iodine. Note the appearance of each sample.

•• Clean the surface of each sample by rubbing them with sand paper and note their appearance again.

•• Try to cut these elements with a sharp knife and note your observations.

•• Hold a piece of sodium with a pair of tongs.

Caution: Always handle sodium with care. Dry it by pressing between the folds of a filter paper.

•• Put it on a watch-glass and try to cut it with a knife.

•• What do you observe?

•• Place any one element on a block of iron and strike it four or five times with a hammer. What do you observe?

•• Repeat above steps with other elements.

•• Record the change in the shape of these elements.

•• Which of the above elements are available in the form of wires?

1.12.1 Properties of Metals

ACTIVITY 1.8: Illustrating Conductivity of Heat and Electricity of Metals

•• Take an aluminium or copper wire. Clamp this wire on a stand, as shown in Figure (a).

Chemical Bonding 15

•• Fix a pin to the free end of the wire using wax.

•• Heat the wire with a spirit lamp, candle or a burner near the place where it is clamped.

•• What do you observe after some time?

•• Repeat the same with carbon or sulphur.

•• Note your observations.

Does the element melt?

Stand	Metal wire
Stand			Free end
Clamp			Free end
			of wire


			Pin
		Burner	Wax

(a)

•• Consider elements aluminium, copper, sulphur and carbon.

•• Set up an electric circuit as shown in Figure (b).

Battery

Bulb

_Clips Switch

A B

Insert sample

to be tested

(b)

•• Place the element to be tested in the circuit between terminals A and B as shown. Does the bulb glow? What does this indicate?

•• Compile your observations regarding properties of elements in your exercise book.

•• Metals are good conductors of heat and electricity. This means that metals allow heat and electricity to pass through them easily. Silver metal is the best conductor of heat. Copper metal is a better conductor of heat than aluminium metal.

•• Metals are lustrous (or shiny). This means that metals have a shiny appearance.

16 S2 Chemistry

•• Metals are usually strong. For example, iron metal (in the form of steel) is very strong when freshly cut and is used in the construction of bridges, buildings and vehicles. Some metals are not strong. For example, sodium and potassium.

•• Metals are ductile. This means that metals can be drawn (or stretched) into thin wires.

•• Gold and silver are among the best ductile metals.

•• Metals are malleable. This means that metals can be hammered into thin sheets.

The cooking utensils are made of metals because metals are good conductors of heat.

EXERCISE 1.11

Name the metal which is the best conductor of electricity.

Aluminium is a better conductor of heat than copper. (True or False)

Metals are _______. This means that they can be hammered into thin sheets.

Why are cooking utensils made of metals?

Which of the following statement(s) is/are correct for metals?

(a) Metals such as sodium and potassium are not strong.

(b) Iron is used in the construction of buildings.

(d) All of these.

1.13 SUMMARY

•• An atom achieves a stable electronic configuration by losing, gaining or sharing electrons.

♦♦Metal atoms with one, two or three electrons in the outermost shell lose electron(s) to form positively charged ions (cations).

♦♦Non-metal atoms with five, six or seven electrons in the outermost shell gain three, two and one electron(s) to form negatively charged ions (anions).

♦♦Non-metal atoms with four to seven outermost electrons may gain electrons by sharing them with each other.

•• A chemical bond is a force that holds ions, molecules or atoms together. A bond is formed when each atom acquires a stable electronic configuration like noble gas.

•• The electrostatic binding force is called an ionic bond or electrovalent bond.

•• Ionic compounds are formed by attraction of positive and negative ions. These compounds are crystalline solid. They conduct electricity. Ionic compounds have high melting and boiling points.

•• A covalent bond forms between two or more atoms of non-metals that are unable to form ions.

Chemical Bonding 17

•• Covalent compound is formed when atoms achieve a stable electronic configuration by sharing of electrons. Covalent compounds are solids, liquid or gases. Covalent compounds have low melting and boiling points.

•• The two forms of carbon that join covalently to form giant structure are diamond and graphite.

•• The force which binds various metal atoms together is called metallic bond.

•• Metals are generally hard, lustrous, strong, malleable and ductile. They conduct heat and electricity in both molten and solid state.

1.14 GLOSSARY

•• Anion: a negatively charged ion.

•• Cation: a positively charged ion.

•• Crystal: a solid where the atoms form a periodic arrangement.

•• Diamond: one of the known allotropes of carbon.

•• Ductile: able to be drawn out into a thin wire.

•• Electronic configuration: the distribution of electrons of an atom.

•• Graphite: a grey crystalline allotropic form of carbon which occurs as a mineral in some rocks.

•• Malleable: able to be hammered or pressed into shape without breaking or cracking.

•• Noble gas: the gaseous elements helium, neon, argon, krypton, xenon, and radon.

1.15 UNIT ASSESSMENT

I. Multiple Choice Questions

The number of electrons gained by non-metals to achieve noble gas electronic

configuration is

(a) one (b) two (c) three (d) all of these

The electronic configuration of sodium ion is

(a) 2,8,1 (b) 2,8,8 (c) 2,8 (d) 2,8,2

The electronic configuration of chloride ion is

(a) 2,8 (b) 2,8,8 (c) 2,8,7 (d) 2,8,3

Choose the ionic compound.

(a)	Calcium chloride		(b)	Copper sulphate
(c)	Sodium hydroxide		(d)	All of these
5. Most ionic compounds are soluble in
(a)	water	(b) petrol	(c)	kerosene	(d) all of these

18 S2 Chemistry

Which of these is not a covalent compound?

(a) Carbon dioxide (b) Methane (c) Ammonia (d) None of these

Choose the correct statement.

(a) Covalent compounds have low melting points

(b) Ionic compounds have high melting points

(d) All of these

Graphite is used for making ____________.

(a) lubricant oils (b) pencil leads (c) both (a) and (b) (d) jewellery

If we burn diamond, the product formed is ____________.

(a) carbon dioxide (b) hydrogen gas

The force which binds various metal atoms together is called ____________.

(a) metallic bond (b) covalent bond (c) ionic bond (d) none of these

Open Ended Questions

How can an atom achieve stability?

Distinguish between covalent and ionic bond.

Compare between the properties of ionic and covalent compounds.

Explain the formation of sodium ion.

Give five examples of each

(a) Ionic compounds (b) Covalent compounds

Compare the conductivity of distilled water with sodium chloride solution.

Write two uses of diamond.

Draw the structure of graphite.

Illustrate the physical properties of metals.

III. Practical-based Questions

Look at the figures and choose the correct statement.

Cl C Cl

H C H

Figure A Figure B

Chemical Bonding 19

(a) Figure A is an example of ionic compound

(b) Figure B is not an example of covalent compound

(d) None of these

The following figure illustrates the electronic configuration of

+11

(a) Lithium (b) Sodium (c) Chlorine (d) Helium

3. shows the structure of ...........................

(a) the hardest substance known (b) an allotrope of carbon

Which of the following materials makes the circuit complete when inserted in between the crocodile clips?

Battery

Bulb

Clips		Switch
Clips
A	B
(a) Aluminium foil	(b)	Copper wire
(c) Both (a) and (b)	(d)	Sulphur
20 S2 Chemistry

In the given figure, arrow shows the

(a) carbon rod (b) iron rod (c) brass rod (d) copper rod

Which of the following depicts the molecule of water?

(a) (b)

PROJECT

Make a 3D model of diamond and graphite and discuss their physical properties.

Chemical Bonding 21

Unit 2

Trends in Properties of Elements in the Periodic Table

LEARNING OBJECTIVES KNOWLEDGE GAIN

After reading this unit, you will be able to:

•• describe trends in reactive elements with acids, water, and halogens.

•• explain the trends in the physical properties across a period and down a group.

Jons Jakob Berzelius was a Swedish chemist and one of the founders of modern chemistry. He proposed the first letter (or first letter and another letter) of the name of an element as its symbol.

2.1 CLASSIFICATION OF ELEMENTS

ACTIVITY 2.1: Distinguishing Metallic and Non-metallic Objects

Collect five objects made of metals. Also collect five objects made of non-metals. Compare the physical properties of metallic objects and non-metallic objects.

Observe Figure 2.1 of the periodic table. There are 118 chemical elements known at present. These elements are classified into metals, metalloids and non-metals. The metals appear at the left-hand side and middle part of the periodic table. The non-metals appear at the right-hand side of the periodic table (Figure 2.1). Metalloids lie in between metals and non-metals.

ACTIVITY 2.2: Categorising Elements into Metals, Non-metals and Metalloids

In groups, classify all elements of periodic table into metals, metalloids and non-metals.

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GSNDA NYUNDO

CHEMISTRY S2

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